1. Fundamentals of General, Organic and Biological Chemistry
7th Edition
Chapter Four
Molecular Compounds

2. Chapter 4 Goals Topics and Goals for the Chapter:
What is a covalent bond?
Be able to define and describe the nature of covalent bonds and how they are formed.
How does the octet rule apply to covalent bond formation?
Be able to use the octet rule to predict the numbers of covalent bonds formed by common main group elements.
Be able to define single bond, double bond and triple bond.
What are the major differences between ionic and molecular bonds?
Be able to compare the structures, compositions, and properties of ionic and molecular compounds.
How are molecular compounds represented?
Be able to interpret molecular formulas and draw Lewis structures for
molecules.
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3. Chapter 4 Goals Cont.
What is the influence of valence-shell electrons on molecular shape?
Be able to define VSEPR theory and be able to apply it to predict the
molecular geometry around a central atom in a Lewis structure.
When are bonds and molecules polar?
Be able to define electronegativity and polar covalent bond.
Be able to understand the trends of increasing and decreasing electronegativity values in the periodic table.
Be able to use electronegativity and molecular geometry to predict bond and molecular polarity.
How are binary molecular compounds named?
Be able to correctly name binary molecular compounds.
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4. 4.1 Covalent Bonds Ionic bonds
Complete transfer of e- from one atom to another to form bond.
Covalent bonds
Sharing of electrons between atoms to form bond.
Generally interactions between nonmetallic elements.
Molecule – a group of atoms held together by covalent bonds.
Nonmetals can achieve an octet by sharing an appropriate number of electrons
in covalent bonds.
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5. H2 is the simplest covalent bond H∙ ∙H 1s 1s 1e- 1e-
When two H atoms come together, electrical interactions occur.
Some interactions are repulsive—the two positively charged nuclei repel each other and the negatively charged electrons repel each other.
Other interactions are attractive—each nucleus attracts the electrons and each electron attracts both nuclei.
When the attractive forces are stronger than the repulsive forces, a covalent bond is formed and the atoms stay together.
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6. Both nuclei attract both e-. Both e- attract both nuclei.
When the nucleus–electron attractions (blue arrows) are greater than the nucleus–nucleus and electron–electron repulsions (red arrows), the result is a net attractive force that holds the atoms together to form a molecule.
Both nuclei attract both e-.
Both e- attract both nuclei.
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7. H-H, H:H and H2 all represent a molecule of hydrogen.
The Hydrogen Molecule
H-H, H:H and H2 all represent a molecule of hydrogen.
Covalent bond in H2 = two spherical 1s atomic orbitals overlap to form molecular
orbital.
Both atoms share the two valence electrons.
The two electrons are concentrated between the two nuclei.
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8. Atoms too close – nuclei repel
In H2 the repulsive and attractive forces depend on how close the atoms are
to one another.
Atoms too close – nuclei repel
Atoms too far away – attractive forces are too small
Optimal distance = minimal energy = bond length = 74 pm for H2
A graph of potential energy versus internuclear distance for the H2 molecule.
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9. Each Cl atom has an octet of electrons.
Another example:
Cl2
Covalent bond in Cl2 = overlap of the two 3p atomic orbitals containing the single unpaired electrons = molecular orbital
Each Cl atom has an octet of electrons.
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10. There are 7 diatomic elements – H2, N2, O2, F2, Cl2, Br2, I2
Atoms are held together by covalent bonds.
N2 and O2 - colorless, odorless, nontoxic gases.
F2 - a pale yellow, highly reactive gas.
Br2 - a dark red, toxic liquid
I2 - a violet, crystalline solid.
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11. 4.2 Covalent Bonds and the Periodic Table
Just looked at molecules formed by atoms of 1 type of element.
Molecules also form through covalent bonding of atoms of different types of
elements.
Examples:
Each atom shares
enough electrons to
achieve a noble gas
configuration.
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12. In general, # of bonds formed follows fulfilling the octet rule.
Figure 5.4 summarizes the number of covalent bonds typically formed by common main group elements.
In general, # of bonds formed follows fulfilling the octet rule.
There are exceptions to the octet rule.
Numbers in parentheses indicate possible numbers of bonds that result in exceptions to the octet rule.
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13. Exceptions to the Octet Rule
Boron
Only has 3e- it can share.
Only forms 3 covalent bonds.
Ex: BF3
Elements in 3rd row and below
Have vacant d orbitals available.
Can have greater than an octet of electrons around the atom.
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14. 4.3 Multiple Covalent Bonds
Atoms can share more than 1 electron.
Double and triple covalent bonds can form between two atoms.
CO2 and N2 are covalent molecules.
How do the atoms obtain a full octet?
Single bonds?
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15. Carbon, nitrogen, and oxygen are the elements most often present in
CO2 must have C, O double
bonds if each atom is to
obtain an octet.
N2 must have a N,N triple
bond if each atom is to
obtain an octet.
Carbon, nitrogen, and oxygen are the elements most often present in
multiple bonds.
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16. Multiple covalent bonds are very common in organic compounds.
Examples:
Ethylene (C2H4) – A compound used commercially to induce ripening in fruit.
Acetylene (C2H2) - The gas used in welding.
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17. 4.4 Coordinate Covalent Bonds
A covalent bond that forms when 1 electron is donated by each atom.
Coordinate covalent bond
A covalent bond that forms when both electrons are donated by the same atom.
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18. 4.5 Characteristics of Molecular Compounds
Molecules have forces between molecules – intermolecular forces
We will look at intermolecular forces more closely in Chapter 8.
If molecules have weak intermolecular forces, the molecules are very weakly attracted to one another.
The molecular substance is a gas at ordinary temperatures.
As intermolecular forces increase, the substances tend to be liquids and solids.
Intermolecular forces weak strong
State gases < liquids < solids
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19. Molecular compounds do not conduct electricity.
The melting point and boiling point of molecular solids are usually lower than those of ionic solids.
Ionic Molecular
NaCl C2H6 (ethane)
m.p. 801ºC ~-200ºC
b.p. 1413ºC ~-150ºC
Molecular compounds do not conduct electricity.
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20. Example of a coordinate covalent bond: Ammonium ion
The filled orbital on N overlaps with the vacant orbital on H+
Once formed, there is no difference between a covalent and a coordinate covalent bond.
Coordinate covalent bonds can result in unusual bonding patterns – i.e. N with 4 bonds rather than 3.
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21. Different Structural Representations for Molecules
4.6 Molecular Formulas and Lewis Structures
Different Structural Representations for Molecules
Molecular formula – A formula that shows the numbers and kinds of atoms in one molecule of a compound.
Condensed structure – Provides minimal structural information, actual bonds not shown but inferred.
Structural formula – Provides minimal structural information, bonds to atoms shown as lines.
Space-filling model – More realistic, shows bonding, relative size of atoms, 3-D structure
Ball and Stick Model – Shows bonding, relative size of atoms, 3-D structure
Lewis Structure - A molecular representation that shows both the connections among atoms and the locations of lone-pair valence electrons.
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22. Examples: C2H6 C2H4O Lewis structure Also structural Ball and stick
formula
Ball and stick
Condensed structure
C2H6
C2H4O
Molecular formula
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23. C – black or grey F – light green O – red Cl – greenish yellow
Molecular
Formula
CH4
NH3
H2O
HCl H2
Color-Code
H – white P – dark blue
C – black or grey F – light green
O – red Cl – greenish yellow
N – blue Br – brownish red
S – yellow I – purple
Figure: 05-05UN03
Title:
Space-filling and ball-and-stick models
Caption:
Space-filling and ball-and-stick models of several compounds are shown.
Notes:
Space-filling models are perhaps more realistic pictures of compounds, but ball-and-stick models allow visualization of bonds as well as atoms.
Keywords:
model, space-filling, ball-and-stick
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24. Examples of Lewis Structures
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25. Ionic Formulas vs. Molecular Formulas
Ionic Formula – a ratio of ions – structure is an extended structure.
Molecular Formula – 1 molecule – actual structure.
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26. 4.7 Drawing Lewis Structures
To draw a Lewis structure you must know how the atoms are connected to one another.
Two approaches are used for drawing Lewis structures.
Approach 1 – Useful for organic molecules
Common bonding patterns are followed by atoms.
C forms four covalent bonds.
N forms three covalent bonds and has one lone pair of electrons.
O forms two covalent bonds and has two lone pairs of electrons.
Halogens form one covalent bond and have three lone pairs of electrons.
H forms one covalent bond.
H and halogens – 1 bond = must be at the ends of structures.
Example: No C-H-C, or C-F-C bonds in this approach.
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27. Step 1- Count the total number of valence electrons
Approach 2 – A general stepwise procedure which works for all molecules
A slightly more detailed handout on the steps is provided in class and available on Blackboard.
Basis of approach – Ignore information regarding which atom donated electrons to form a bond. Focus on the total number of atoms and the total number of valence electrons.
Step 1- Count the total number of valence electrons
Molecule – Add valence electrons for all atoms
Polyanion – Add an electron for each negative charge
Polycation – Subtract an electron for each positive charge
Step 2- Draw a skeleton structure for the species, joining atoms by single lines to represent covalent bonds
For some species there might be more than one possible structure if you are given
a formula.
Most of the molecules we deal with have a central atom bonded to two or more terminal atoms (NH4+, SO2, CCl4).
The central atom is usually written first in the formula followed by terminal
atoms.
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28. Step 4- Place all remaining electron lone pairs on the central atom.
Approach 2 – A general stepwise procedure which works for all molecules - Continued
Step 3- Determine the number of valence electrons still available for distribution and add electrons so that each peripheral atom connected to the central atom gets an octet. (H = 2, not 8 electrons)
That is, deduct two valence electrons for each single bond written in Step 2. and place the remaining electrons around the peripheral atoms in the structure.
Step 4- Place all remaining electron lone pairs on the central atom.
Step 5- If the central atom does not yet have an octet after all electrons have been assigned, take a lone pair from a neighboring atom and form a multiple bond to the central atom.
Note: This will NOT change the number of electrons counted towards the octet on the peripheral atom. You are simply “sharing” the two electrons.
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29. Practicing drawing Lewis structures: O3 – ozone NO3-
When determining the Lewis structure for a molecule – Approach 2 will ALWAYS give you the correct structure.
Practicing drawing Lewis structures:
O3 – ozone NO3-
Try CO2 on your own.
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30. 4.8 The Shapes of Molecules
If we look at 3-D drawings of molecules you notice each molecule has a specific shape.
Bent
Pyramid-Shaped
Tetrahedral
Molecular shapes are related to number and location of valence e- around the atoms.
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31. We can predict molecular shapes! Use VSEPR Model
Valence Shell Electron Pair Repulsion Model
According to VSEPR:
Valence e- pairs (in bonds and unshared) surrounding an atom repel each other.
Bonds and unshared e- orient themselves as far away from one another as possible to minimize repulsion.
They take advantage of all available space.
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32. 3 steps to apply model to determine molecular shape of a molecule:
VSEPR Model
3 steps to apply model to determine molecular shape of a molecule:
Step 1: Draw a Lewis structure of the molecule and identify the atom whose geometry is of interest.
In a simple molecule – The atom of interest will be the central atom.
Step 2: Count the number of electron charge clouds (areas of electron density) surrounding the atom of interest.
Single bond = 1 e- cloud
Double bond = 1 e- cloud
Triple bond = 1 e- cloud
Lone e- pair = 1 e- cloud
Step 3: Predict molecular shape by assuming that the charge clouds orient in space so that they are as far away from one another as possible.
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33. Focusing on central atom
Solid line – bond in plane of
page
Solid wedge – bond coming
out of plane of page
Dashed line – bond projecting
into the plane of the page
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34. Linear molecules have bond angles of 180°
Linear molecules have bond angles of 180°. Planar triangular molecules have bond angles of 120°. Tetrahedral molecules have bond angles of 109.5°.
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35. Molecular geometry of CH4, NH3 and H2O
Lone e- pairs are more repulsive:
2 lone e- pairs > lone e- pair > 0 lone e- pairs
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36. Look at molecule as having several central atoms.
Geometry around atoms in larger molecules is also derived from shapes in Table 5.1.
Ex:
Look at molecule as having several
central atoms.
Planar
triangular
120º
Tetrahedral
109.5º
Bent
104.5º
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37. What is the molecular geometry of NO3-?
Question:
What is the molecular geometry of NO3-?
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38. 4.9 Polar Covalent Bonds and Electronegativity
Electrons in a covalent bond occupy the region between bonded atoms.
Bond polarity – Describes sharing of e- between atoms.
Non-polar covalent bond – Electrons shared equally between two atoms.
Atoms need to be identical
Ex: Cl2, H2, N2, O2
Polar covalent bond – Electrons shared unequally between two atoms.
Atoms are not identical
Bonding e- are attracted more strongly by one atom than by the
other.
Ex: H-O-H and H-Cl
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39. Molecules with Polar Covalent Bonds
Molecules are neutral overall but they have + and – areas.
d+ = partial positive placed on more positive atom
d- = partial negative placed on more negative atom
H-O-H H-Cl d+ d- d+ d+ d-
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40. Electrostatic potential map – Uses color to portray e- distribution
A great way of looking at unequal distribution of bonding electrons is to look at a electrostatic potential map.
Electrostatic potential map – Uses color to portray e- distribution
Blue = + red-yellow = -
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41. Electronegativity increases across a period.
We can determine bond polarities and predict the occurrence of covalent vs. ionic bonding in a molecule by comparing electronegativities of atoms.
Electronegativity – The ability of an atom in a molecule to attract e- to itself.
The electronegativity scale is based on giving the most electronegative element,
F, an assigned value of 4.
Electronegativity increases across a period.
Electronegativity decreases down a group.
Noble gases are not assigned values.
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42. Predicting bond polarity
Compare electronegativities of atoms in bonds:
H―C C―N C―O H―Cl Na Cl
E.N
Partial Charge d+ d- d+ d- d+ d d+ d- d+ d-
E.N difference
Rule of thumb:
E.N. difference < Only slightly polar covalent bonds (H―C)
0.5 < E.N. difference < Increasingly polar bonds (C―N, C―O, H―Cl)
2.0 ≤ E.N. difference Ionic bonds (Na+ Cl-)
As E.N. difference increases, bond polarity increases
nonpolar covalent (0) < slightly polar < very polar < ionic
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43. Sample problem:
Predict whether the H-S bond in H2S is nonpolar covalent, polar covalent or ionic.
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44. 4.10 Polar Molecules
Just as we can describe bonds as being polar or nonpolar, we can describe molecules as being polar or nonpolar.
Polarity of molecules – Due to the sum of all individual bond polarities and lone pair contributions in the molecule.
The physical properties of a molecule are dependent on the polarity of the molecule.
Dipoles or polarity can be represented by an arrow pointing to the negative end of the molecule with a cross at the positive end resembling a + sign. d+ d-
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45. Electrostatic potential maps show polarity of molecules very clearly.
Molecular polarity depends on the shape of the molecule as well as the bonds.
Electrostatic potential maps show polarity of molecules very clearly.
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46. Polarity of as molecule depends on the shape of the molecule.
Be careful! Not all molecules with polar covalent bonds are polar overall.
Polarity of as molecule depends on the shape of the molecule.
Symmetrical shapes cause the individual bond polarities to cancel each other out.
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47. A few practice problems:
Determine whether the following molecules are polar:
BH3
SO32-
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Chapter 4

48. 4.11 Naming Binary Molecular Compounds
Binary compounds – Formulas of molecular compounds with two different types of
atoms.
Rules:
1. Formulas are written with the least electronegative element first
a. Metals always written before nonmetals.
b. Nonmetals further to the left on the periodic table are written first.
N before O
2. When naming you must identify exactly how many atoms of each element are in the molecular formula.
3. Prefixes are used to indicate the number of atoms of each element that combine.
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49. Naming Binary Molecular Compounds Rules Cont.
4. Two steps for naming:
a. Name the first element in the formula.
b. Name the second element in the formula.
Use an –ide ending as for anions.
For both of the above steps use a prefix
to indicate the number of atoms,
if needed.
Note: The prefix mono can be omitted
for the first element. (It is assumed to
be one if given no prefix.)
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50. H2O2 dihydrogen dioxide (common name hydrogen peroxide)
Examples:
H2O2 dihydrogen dioxide (common name hydrogen peroxide)
CS2 carbon disulfide
CF4?
NO?
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51. Putting It All Together
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52. Chapter Summary
A covalent bond is formed by the sharing of electrons between atoms rather than by the complete transfer of electrons from one atom to another.
Two shared electrons are a single bond , four are a double bond, and six are a triple bond.
The group of atoms held together by covalent bonds is called a molecule.
When a lone pair of electrons on one atom overlaps a vacant orbital on another atom a coordinate covalent bond is formed.
An atom shares enough electrons to reach a noble gas configuration.
Molecular formulas show the numbers and kinds of atoms in a molecule.
Lewis structures show how atoms are connected in molecules.
Covalent bonds are indicated as lines between atoms, and valence electron lone pairs are shown as dots.
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53. Chapter Summary Cont.
Molecules have shapes that can be predicted using the VSEPR model.
The electronic geometry of atoms with 2 electron charge clouds is linear, with 3 it is planar triangular, and with 4 it is tetrahedral.
Bonds between atoms are polar if the bonding electrons are not shared equally between the atoms.
The ability of an atom to attract electrons is electronegativity. It is highest on the upper right of the periodic table and lowest on the lower left.
Molecular polarity is the sum of all individual bond polarities and lone pair contributions in a molecule.
Molecular compounds usually have lower melting points and boiling points than ionic compounds, many are water insoluble, and they do not conduct electricity when melted or dissolved.
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